Acid-Base Reactions and Redox PDF

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Acid-Base Reactions and Redox

Properties of Acids and Bases Property Taste Litmus paper Phenolphthalein

Sour Red Colorless

Acids

Bases Bitter, Chalky Blue Pink

Touch Arrhenius definition

Corrosive Produces H+ ions in water

Feels slippery, soapy Produces OH- ions in water

For example, carbonic acid is formed when the carbonate ion (CO32-) bonds with two hydrogen ions (H+) to give H2CO3. Other common acids are listed below: Hydrochloric acid: HCl Sulfuric acid: H2SO4 Nitric Acid: HNO3 Acetic Acid: HC2H3O2

Rules for Naming Acids Do Not Contain Oxygen in the Anion:   

Since all these acids have the same cation, H+, we don't need to name the cation. The acid name comes from the root name of the anion name. The prefix hydro- and the suffix -ic are then added to the root name of the anion. HCl, which contains the anion chloride, is called hydrochloric acid. HCN, which contains the anion cyanide, is called hydrocyanic acid.

Naming Oxyacids (anion contains the element oxygen):   

Since all these acids have the same cation, H+, we don't need to name the cation. The acid name comes from the root name of the oxyanion name or the central element of the oxyanion. Suffixes are used based on the ending of the original name of the oxyanion. If the name of the polyatomic anion ended with -ate, change it to -ic for the acid and if it ended with -ite, change it to -ous in the acid. HNO3, which contains the polyatomic ion nitrate, is called nitric acid. HNO2, which contains the polyatomic ion nitrite, is called nitrous acid.

Strong Acids and Bases Acids and bases are called “strong” if they dissociate completely (or almost completely) in water. For example, hydrochloric acid (HCl) exists almost completely as H+ ions and Cl- ions in water. Therefore HCl is a strong acid. If you put 2.0 moles of HCl in water, the HCl will dissociate and you will end up having about 2.0 moles of H+ and 2.0 moles of Cl- in the water. Each HCl molecule breaks up into two ions when dissolved: HCl  H+ + Cl-. For every mole of HCl added to water, you will get one mole of H+.

Weak Acids and Bases Weak acids and bases do not dissociate completely in water. For example, consider acetic acid (HC2H3O2): HC2H3O2  H+ + C2H3O2For HCl, the concentration of HCl equaled the concentration of H+ so finding the pH was easy. However, for acetic acid, [HC2H3O2] does not equal [H+] because not all of the HC2H3O2 dissociates into H+ ions.

Neutralization Reactions Acid- Bases reactions are called neutralization reactions where a salt (ionic compound) and water are formed. For example: Hydrochloric acid reacting with Sodium hydroxide HCl (aq) + NaOH(aq)  H2O(l) + NaCl(aq) Acid + Base  water + Salt

Redox Reactions An example of redox reaction is when magnesium reacts with oxygen to produce magnesium oxide by a transfer of electrons. 2 Mg (s) + O2(g)  2 MgO(s) To illustrate how the electrons move, we write half reactions that show the electrons in the equation. Loss = oxidation 2 Mg  2 Mg2+ + 4 eO2(g) + 4 e-  2 O2Gain = reduction Some redox reactions occur without a change in the number of valence electrons. To show this the concept of oxidation number was developed and new definition where created for oxidation and reduction. Oxidation involves an increase in the oxidation number of an atom. Reduction occurs when the oxidation number of an atom decreases. By definition, the oxidation number of an atom is equal to the charge that would be present on the atom if the compound was composed of ions. The key to identifying oxidation-reduction reactions is recognizing when a chemical reaction leads to a change in oxidation number for one or more atoms.

Oxidation Rules Rule The oxidation state of an atom in an uncombined element is zero. The oxidation state of a monatomic ion is the same as its charge. Oxygen is assigned as oxidation state of -2 in most of its covalent compounds. Hydrogen oxidation state is +1

Most electronegative nonmetals in compounds are assigned oxidation states based on their positions on the Periodic Table. Group # - 8 For a neutral compound the sum of the oxidation states = zero For a polyatomic ion, the sum of the oxidation states = charge

Examples Na : 0 O2 : 0 Na+ : +1 Fe3+ : +3 In CO : -2 In SO2: each oxygen -2 Exceptions: O22- (peroxide) : -1 In NH3: each hydrogen +1 Exceptions: MHx where M is a metal then hydrogen -1 In NaF: fluoride -1 (Group 7-8)

In NaCl: Sodium +1, chloride -1 In CO32Each oxygen = -2 total of -6 Carbon = +4 So that when you add them together you get the charge of 2-

Half-Reactions -

Assign oxidation numbers (they will be placed above each element in parentheses) Split equation into two half reactions For each half reaction o Balance elements o Balance charges by adding electrons  Note: total charge the reactants = total charge of product

(0) (0) (-1) (+3) Example 1: 3 Cl2 + 2 Al 6 Cl- + 2 Al3+ Each chlorine changes from a 0 to a -1 oxidation state  reduction -Electrons need to be placed on the reactant side to show a gain. -There are a total of six chlorine atoms so there is a gain of 6 e6e- + 3 Cl2 Total Charge: 6(-1)= 6 (-1) -6 = -6

6 Cl-

Each aluminum changes from a 0 to a +3 oxidation state  oxidation -Electrons need to be placed on the product side to show a loss. -There are a total of six aluminum atoms so there is a loss of 6 e2 Al 2 Al3+ + 6eTotal Charge: 0 = 2(3+) + 6(-1) 1 =0

An oxidizing agent is a reactant molecule that causes oxidation and contains the element reduced. A reducing agent is a reactant molecule that causes reduction and contains the element oxidized. Example 2:

(+2) (0) (0) (+2) 2+ Cu + Mg Cu + Mg2+ Each copper atom changes from +2 to 0 oxidation state  reduction -Electrons need to be placed on the reactant side to show a gain. -There is only one copper atom so there is a gain of 2 e2e- + Cu2+ Cu Total Charge: 2(-1) + 1(2+)= 0 1 =0

Each magnesium atom changes from 0 to +2 oxidation state  product -Electrons need to be placed on the product side to show a gain. -There is only one magnesium atom so there is a loss of 2 eMg Mg2+ + 2eTotal Charge: 0 = 1(2+) + 2(-1) 1 =0

Species oxidized: Magnesium Oxidizing agent: Cu2+ Species reduced: Copper Reducing agent: Mg As the reaction proceeds, electrons are transferred from Mg to Cu2+...